Introduction to energetics and enthalpy

Introduction to energetics and enthalpy


In this lesson, we will learn:

  • How energy changes are measured in chemical reactions.
  • The standard conditions at which enthalpy changes are reported.
  • To construct enthalpy diagrams for endothermic and exothermic reactions.
  • The definitions of enthalpy of reaction, formation, combustion and neutralization.


  • Outside of chemistry, to be energetic means to have energy.
    In chemistry, ‘energetics’ is about changes in amount of energy in chemical reactions, going from reactants to products. This is very useful to chemists because many chemical reactions give out or absorb energy as heat that we can measure.
    The net amount of heat energy absorbed or released from a reaction at constant pressure, is called an enthalpy change, given the symbol ΔH\Delta H (‘delta H’).
    • Measuring and reporting enthalpy changes is done at standard conditions. This is to have a reference where different data sets can be easily compared.
      Standard conditions define pressure at 100kPa and 1M concentration for solutions.
      Temperature is not actually part of standard conditions, but for fair comparison, most thermodynamic data collected is at 298 K (25°C).
      This is not the same as standard temperature and pressure (STP), which does strictly define 100 kPa and 273 K!
    • The standard enthalpy change of reaction, H \triangle H^{\ominus} r, is the enthalpy change measured when a reaction takes place at standard conditions with all reactants and products in their standard states. Standard states are the states that you find substances in at standard conditions (such as H2O as liquid water, not as steam/ice).

  • The enthalpy change of reaction comes from two sources. One way of thinking about energy is it being in one of two places: inside the molecules you are studying (the system), or outside the system, in the rest of the universe (the surroundings).
    • To begin a chemical reaction, heat energy must be supplied by the surroundings to break bonds in the reactant molecules. The stronger the bond, the more energy is needed.
      • For example, the N≡N triple bond is a lot stronger than the Cl-Cl bond, so more energy is needed to break the N≡N bond than the Cl-Cl bond.
    • When new bonds form to make the products, energy is released back into the surroundings. The more stable/stronger the bond, the more energy is released when it forms.
      • For example, the strong N≡N triple bond releases a lot more energy when it forms as a product of a reaction than when the Cl-Cl bond forms. Forming the strong N≡N bond releases a lot of energy into the environment, so the N≡N molecule itself is in a lower energy state.
      • This is why chemists sometimes call stable molecules ‘low energy’ – they have released most of ‘their energy’ to the environment. In the same way, unstable molecules are sometimes called ‘high energy’.

  • It’s very rare for a reaction to need a certain amount of energy to break up the reactants, and then give back the exact same amount of energy when products form.
    Almost all chemical reactions have an overall effect of absorbing energy from or releasing energy into the surroundings. There are two definitions for this:
    • An exothermic reaction is a chemical reaction that has the overall effect of releasing heat energy to the surroundings.
      • This means that the energy released when forming product bonds was greater than the energy required when breaking up the reactant bonds.
      • Exothermic enthalpy changes are always given a negative sign, to show that the chemical substances have become lower energy (the energy they released is now in the surroundings).
    • An endothermic reaction is a reaction that has the overall effect of absorbing heat energy from the surroundings.
      • This means that the energy required to break up the reactant molecules was greater than the energy released when the new product bonds formed.
      • Endothermic enthalpy changes are always given a positive sign, to show that the chemical substances have become higher energy (because they absorbed more energy from the surroundings as reactants than they released when they became products).

    These two types of reactions can be shown by drawing an enthalpy level diagram.

  • There are some specific enthalpy definitions that should be known for common chemical changes:
    • The standard enthalpy of combustion, H \triangle H^{\ominus} c, is the enthalpy change measured when one mole of a substance is completely reacted with oxygen at standard conditions.
    • The standard enthalpy of neutralization, H \triangle H^{\ominus} n, is the enthalpy change measured when an acid and base react to form one mole of water in a neutralization reaction at standard conditions.
    • The standard enthalpy of formation, H \triangle H^{\ominus} f, is the enthalpy change when one mole of a substance is formed from its constituent elements in their standard states at standard conditions.
    These definitions have several things in common, which is exactly the point.
    • The definitions are all measuring energy change for one mole of a substance.
    • The definitions are all measurements at standard conditions and substances in their standard states.
    This is all done so scientists can compare different data sets properly. If the same temperature (298K), same pressure (100kPa) and same amount of substance (1 mole) are all used, then we can prove that it is the different substances that produce the different enthalpy changes.
  • Introduction
    Introduction to energetics
    How does energy change in a chemical reaction?

    Exothermic and endothermic reactions.

    Important enthalpy definitions.